Sigma and Pi Bonds (A-Level Chemistry)

Summary

This informative video from chemistrystudent.com dives into the intricacies of sigma and pi bonds, essential concepts in understanding covalent bonding. It explains that covalent bonds form when two atoms share a pair of electrons, with the bond's strength and properties depending on the type of overlap between atomic orbitals. Sigma bonds result from a direct overlap, positioning the bonding orbital close to both nuclei for a strong, non-rotatable bond. In contrast, pi bonds arise from the sideways overlap of p orbitals, creating a weaker, less stable bond with restricted rotation. The video uses ethane and ethene as models to illustrate single and double bonds, respectively, highlighting the increased reactivity of double-bonded molecules like ethene due to the presence of pi bonds. The content is enriched with visual explanations and is aimed at viewers with a foundational understanding of chemistry.

Takeaways

• 🔬 Covalent bonds are formed when two atoms share a pair of electrons, resulting from the overlap of half-filled atomic orbitals from each atom.
• 🧲 The attraction between the positively charged nuclei and the shared electrons in the new orbital pulls the atoms together, forming a covalent bond.
• 📍 Sigma bonds are created from the direct overlap of atomic orbitals, resulting in a strong bond that is close to the nuclei of both atoms.
• 📏 Pi bonds form from the sideways overlap of p-shaped orbitals, creating a weaker bond that is further from the atomic nuclei compared to sigma bonds.
• 🚫 Pi bonds restrict rotation due to the overlapping p orbitals, which can twist and break if rotated too far.
• 🔄 The formation of a sigma bond from one pair of half-filled p orbitals leaves the remaining p orbitals free to potentially form a pi bond.
• 🛑 Sigma bonds are always stronger than pi bonds, which contributes to the differing reactivity between molecules with single and double bonds.
• ⚙️ In ethane, carbon atoms are connected by a sigma bond, with each carbon atom having three bonds to hydrogen atoms.
• ⚓️ In ethene, a carbon-carbon double bond consists of one sigma bond and one pi bond, represented by two lines between the carbon atoms.
• 🔗 A double bond is stronger and less likely to break than a single bond, but a carbon-carbon double bond is not twice as strong as a single bond due to the weaker pi bond.
• 🌟 Alkenes, like ethene, are more reactive than alkanes, such as ethane, because they contain pi bonds that are easier to break, leading to increased reactivity.

Q & A

What is a covalent bond?

A covalent bond is a type of atomic bond formed when two atoms share a pair of electrons. It occurs when half-filled atomic orbitals from two different atoms overlap, creating a bonding orbital where the pair of electrons can exist, attracting both positively charged nuclei to the electron density and pulling the atoms together.

What are the different shapes of atomic orbitals that can be involved in covalent bonding?

The different shapes of atomic orbitals that can be involved in covalent bonding include s and p shaped orbitals, which are the most commonly studied. Other half-filled atomic orbitals can also overlap or merge to create a bonding orbital between two atoms.

How does a sigma bond form?

A sigma bond forms when two atomic orbitals face each other and overlap easily. The resulting bonding orbital is close to the nuclei of both atoms, with both nuclei having a high level of attraction to the electrons in the orbital, making the bond very strong.

What is a pi bond and how does it differ from a sigma bond?

A pi bond is a covalent bond that forms from the sideways overlap or bending of p-shaped atomic orbitals from two atoms. Electrons in a pi bond are further from the nuclei of both atoms compared to a sigma bond, resulting in a weaker attraction and a bond that is easier to break. Unlike sigma bonds, pi bonds have restricted rotation due to the two areas of electron density above and below the sigma bond.

How does the bonding in ethane (C2H6) differ from that in ethene (C2H4)?

In ethane, carbon atoms are bonded by a sigma bond formed by the direct overlap of two orbitals, with each carbon atom having three bonds to hydrogen atoms. In ethene, the carbon atoms are bonded by both a sigma bond and a pi bond, forming a carbon-carbon double bond. Each carbon atom in ethene has two bonds to hydrogen atoms, with two half-filled p-shaped orbitals left over for the pi bond formation.

Why are double bonds stronger than single bonds?

Double bonds are stronger than single bonds because they involve both a sigma bond and a pi bond between the atoms. The presence of two bonds results in a higher bond energy and a stronger attraction between the atoms, making the double bond harder to break than a single bond.

Why are alkenes more reactive than alkanes?

Alkenes are more reactive than alkanes because they contain a carbon-carbon double bond, which includes a pi bond that is weaker and more easily broken than a sigma bond. This higher reactivity allows alkenes to undergo a wider range of chemical reactions compared to alkanes.

What happens if a pi bond is broken?

If a pi bond is broken, the sigma bond between the atoms still remains intact. This leaves the molecule with a single bond between the previously double-bonded atoms, which can lead to different chemical reactivity and properties compared to when the pi bond was intact.

What is the significance of the orientation of p orbitals in forming sigma and pi bonds?

The orientation of p orbitals is significant because it determines the type of bond that can form. When p orbitals overlap directly, they form a sigma bond. However, when they cannot overlap directly due to their orientation, they can bend and overlap sideways to form a pi bond. The three p orbitals (px, py, pz) are oriented at 90 degrees to each other, which allows for the formation of pi bonds after a sigma bond has been established.

How does the rotation of atoms affect the pi bond?

The rotation of atoms affects the pi bond because the p orbitals that form the pi bond overlap sideways. If one atom rotates, the pi bond's orbitals will also attempt to rotate, which can lead to twisting. If twisted too far, the pi bond can break due to the strain caused by the misalignment of the p orbitals.

What is the role of electron density in the strength of sigma and pi bonds?

Electron density plays a crucial role in the strength of sigma and pi bonds. In sigma bonds, the electron density is closer to the nuclei of both atoms, resulting in a strong attraction and a robust bond. In pi bonds, the electron density is further from the nuclei, leading to a weaker attraction and a less stable bond compared to sigma bonds.

Why are sigma bonds free to rotate?

Sigma bonds are free to rotate because the overlapping orbitals that form the bond are directly aligned with the nuclei of the bonding atoms. This direct alignment allows for rotation without disrupting the orbital overlap, which means the bond's strength and integrity are maintained during rotation.